Polarity of Bonds and Molecules

Bond Polarity

Polarity in organic chemistry refers to a separation of charge and can describe a bond
or an entire molecule. Experimentally, bond polarity is measured by its dipole moment.
Bonds connecting atoms of different electronegativity are polar with a higher density of
bonding electrons around the more electronegative atom giving it a partial negative charge
(designated as d-). The less electronegative atom has some
of its electron density taken away giving it a partial positive charge (d+).

This polarization of charge in the H-Cl bond is due to different electronegativities of
chlorine and hydrogen.

The polarity in the bond can also be represented by a arrow indicating a dipole (two
charges separated by a distance). The tip of the arrow points toward the more
electronegative atom.

Molecular Polarity

The polarity of the molecule is the sum of all of the bond polarities in the molecule.
Since the dipole moment (m, measured in Debyes (D)) is a vector
(a quantitiy with both magnitude and direction), the molecular dipole moment is the vector
sum of the individual dipole moments. If we compare the molecular dipole moments of
formaldehyde and carbon dioxide, both containing a polar carbonyl (C=O) group, we find
that formaldehyde is highly polar while carbon dioxide is nonpolar . Since CO2
is a linear molecule, the dipoles cancel each other.

Water is a bent molecule with polar O-H bonds. The bond dipole moments add to give a
resultant dipole (m = 1.85 D) directed toward the more
electronegative oxygen.

The polarity of chloromethanes reveals the importance of symmetry. All of these
compounds contain polar C-Cl bonds but the tetrahedral symmetry of CCl4 causes
the bond dipoles to cancel giving a nonpolar molecule.

Chloromethane

The top image show the bond electron density and the
bottom image the molecular dipole.

m = 1.87 D

Dichloromethane

The top image show the bond electron density and the
bottom image the molecular dipole

m = 1.54 D

Trichloromethane

The top image show the bond electron density and the
bottom image the molecular dipole

m = 1.02 D

Tetrachloromethane

The top image show the bond electron density and the
bottom image the molecular dipole

m = 0 D

Dipoles and Intermolecular Attraction

Melting points and boiling points are important physical properties. These properties
reveal something about the forces that hold molecules together in condensed phases
(liquids and solids). Chemists recognize three major kinds of attractive forces in
covalent molecules, all of which are related to dipoles.

Polar molecules have a permanent dipole moment. Since opposite charges attract, when
polar molecules approach each other they orient themselves in a head-to-tail manner. The
following example shows the dipole-dipole attraction in chloroform (trichloromethane, bp
61oC).

Carbon tetrachloride (tetrachloromethane, bp 77oC) is a nonpolar molecule
but it is a liquid at room temperature, indicating that some attraction between molecules
must exist. The molecule has no permanent dipole but an instantaneous dipole is formed
when two CCl4 molecules approach each other. The electron cloud in one molecule repulses
the electrons in the second molecule breaking the symmetry. These temporary dipoles exist
for only a short time and fluctuate from one molecule to another. The result is a weak
dipole-dipole attraction called the London dispersive force (van der Waals force). The
more contact area between molecules the stronger the van der Waals forces. We will see
examples of this trend when we examine the boiling points of hydrocarbons.

Tetrachloromethane
molecules far apart. No dipole moment.

Induced dipole moment of
two tetrachloromethane molecules close together.

Hydrogen bonding is the result of strong dipole-dipole attraction in molecules
containing O-H and N-H bonds. HF also undergoes hydrogen bonding but since F is
monovalent, this bond is not found in organic molecules. These hydrogen bonds are strongly
polarized with the hydrogen atom carrying a partial positive charge. Although hydrogen
bonding is only about 10% as strong as covalent bonds it is responsible for the unusual
high boiling points of water and alcohols which contain O-H bonds. Ammonia and amines
contain N-H bonds which are less polar than O-H bonds and the resulting hydrogen bonding
is weaker.